Sciencemadness Discussion Board - Separating MgCl2 from a solution of MgCl2 and CaCl2

Separating MgCl2 from a solution of MgCl2 and CaCl2

Eli25 - 11-5-2015 at 01:11

I have a solution containing MgCl2 and CaCl2 with water. I need to separate MgCl2.How can I do it in the best and easiest way?

annaandherdad - 11-5-2015 at 05:11

Apart from fractional crystallization, you could add sodium sulfate to precipitate the calcium, since magnesium sulfate is much more soluble in water than calcium sulfate. I'm not sure this will work, however, since the solubility of calcium sulfate is affected by the presence of the chlorides. In any case, calcium sulfate is still slightly soluble in water, and even with the calcium gone, you'd end up with a mixture of sodium chloride and magnesium sulfate.

Alternatively, you could precipitate both calcium and magnesium as carbonates, then dissolve the (mostly magnesium carbonate) in sulfuric acid to get magnesium sulfate. This could be later converted to the chloride, by going through the carbonate again, if necessary. Both magnesium and calcium carbonates are quite insoluble in water.

I've been working on separating sea water into its ingredients, just for fun. At a certain point you get a mixture of magnesium and calcium chlorides (mostly). I'm using fractional crystallization. There is a very interesting thread here on SM on separating sodium and potassium chlorides, you should check it out.

DistractionGrating - 11-5-2015 at 12:02

Wouldn't another option be to add NaOH or KOH to precipitate Mg(OH)2, and then add HCl to the Mg(OH)2?

gboneu - 12-5-2015 at 11:17

That wouldn't work because both Mn(OH)2 and Ca(OH)2 are insoluble in water, and will both react with the acid to reform the chlorides.
The method i have in my book is the next.
If you have Calcium contamination un your Manganese chloride, precipitate all the manganese with Hydrogen sulfide (SH2) until no more comes out, then filter off and wash with water to wash the soluble contaminants and acids. Then dissolve all the manganese sulfide in hydrochloric acid and heat up the solution to get out all the SH2 remaining in solution.
Obviously a recristalization is recomended

If you want to see if there are still some calcium ions you can test them with these reagents:
Get a small drop and touch a filter papaer, leave it to dry and test...
1)SODIUM RHODIZONATE (Gray-blue colour with Ca)
_Sodium Rhodizonate (1% Aqueous solution ) Spray 1
_ Ammonia (25% aqueous solution) Spray 2
2)ALIZARINE (Red-Violett colour with Ca)
_ Alizarine in Chloroform (2% solution) Spray 1
_2M sodium hydroxide solution. Spray 2_ Ammonia solution (25%) Spray 3
Note. If you dont have chloroform usea Alcoholic alizarine solution (saturated)
3) QUINALIZARIN (Violet colour with Ca)
_Quinalizarin solution (0.05%) in Ethanol (70% by volume) Spray 1
_Ammonia solution (25% aqueous solution) Spray 2
NOTE: If you have negative resoults try putting the papers in a ammonia rich atmosphere, the original analisis is using ammonia gas, bur the solution has to work alsow, but it's always better to know

DistractionGrating - 12-5-2015 at 11:40

Gboneu, I think you are confusing Manganese (Mn) with Magnesium (Mg).

You have a point, though, that Ca(OH)2 is not very soluble in water, but significantly more so than Mg(OH)2. I guess my idea's feasibility depends on how much Ca is in the solution in question. I got my idea from the Hach Calcium Hardness EDTA Titration method, where KOH is used to precipitate the Magnesium, leaving Calcium in solution.

Chemosynthesis - 12-5-2015 at 12:04

Coulometry.

gboneu - 12-5-2015 at 13:14

Owww.. (-.-)

gboneu - 12-5-2015 at 13:20

Ok, i found some another info in my book's, Magnesium carbonate is soluble in a solution of ammonium chloride and ammonia, where calcium strontium and barium are not soluble (There carbonates) so you can extract the magnesium with this solution, Then maybe ammonium carbonate can be decomposed by hydrochloric acid i think ... Sorry for the confusion, im was exited when i found that info xD

[Edited on 12-5-2015 by gboneu]

blogfast25 - 12-5-2015 at 14:04

Quote: Originally posted by gboneu

Ok, i found some another info in my book's, Magnesium carbonate is soluble in a solution of ammonium chloride and ammonia, where calcium strontium and barium are not soluble (There carbonates) so you can extract the magnesium with this solution, Then maybe ammonium carbonate can be decomposed by hydrochloric acid i think ... Sorry for the confusion, im was exited when i found that info xD

I'm not convinced that this information is completely correct.

Magnesium hydroxide is indeed soluble in ammonium salt solutions due to:

Mg(OH)2 + 2 NH4+ === > Mg2+ + 2 NH3 + 2 H2O

Magnesium carbonate, acc. Holleman's 'Inorganic Chemistry', is a basic carbonate: Mg(OH)2.3MgCO3.3H2O. Whether the ammonium salt trick also works on that basic carbonate will depend on its solubility product Ks.

gboneu - 12-5-2015 at 14:13

Here says that it work's with the carbonate.

blogfast25 - 12-5-2015 at 17:52

Quote: Originally posted by gboneu

Here says that it work's with the carbonate.

Forgot to provide the link, maybe? I'm quite interested in seeing it. Ta.

[Edited on 13-5-2015 by blogfast25]

Turns out there are several Mg carbonates (Wiki), including a 'straight' MgCO3, with a Ks of 3.5×10–8. That's probably soluble enough to soluble in ammonium salts...

[Edited on 13-5-2015 by blogfast25]

gboneu - 13-5-2015 at 00:46

The book where i got the info is the (I live in argentina it's in spanish)
https://books.google.com.ar/books?id=LoTwmQEACAAJ&dq=Pra...

gboneu - 13-5-2015 at 00:54

I search some info on using Chromatography but it's a little complicated to do, And it takes a lot of time, However if i have time i give it a try on my lab to se what can i get.

Eli25 - 15-5-2015 at 09:10

Hi guys. Thanks for your answers. Actually I want MgCl2 as the final product ...I just need to remove Ca from solution.....pleaseeeeee help me... I need to keep my job by giving the best solution...

What if I add MgSo4 and remove CaSo4 as precipitate ? Can I have pure MgCl2 solution?

I'm working in a factory so the answer should be practical and economical....

blogfast25 - 15-5-2015 at 09:45

Quote: Originally posted by Eli25

Hi guys. Thanks for your answers. Actually I want MgCl2 as the final product ...I just need to remove Ca from solution.....pleaseeeeee help me... I need to keep my job by giving the best solution...

What if I add MgSo4 and remove CaSo4 as precipitate ? Can I have pure MgCl2 solution?

Yes but CaSO4 is better described as 'poorly soluble' than 'insoluble'. Its solubility (acc. Wiki) is about 2 g/litre, so some residual CaSO4 will stay in solution. It depends whether this relatively small contamination is acceptable to you or not.

You'd also need to know the concentration of CaCl2 quite well, as otherwise you may leave more Ca in solution or contaminate the MgCl2 with excess MgSO4.

[Edited on 15-5-2015 by blogfast25]

Chemosynthesis - 15-5-2015 at 09:59

I am not sure how practical this is for you, but you may wish to consider trying a calcium specific ionophore such as ETH 129, ETH 5234, or ETH 1001. These work in coulometric separations, and given the lipophilicity, could probably be utilized in a solvent extraction with something you have on hand, though the latter may not be an option for you due to purity standards and the addition of an organic process stream.

Edit: and, probably sounding extremely callous here, but while I am sorry your job is at risk, that isn't really our problem or what the forum is for. I wish you all the best, but if your job retention comes down to what strangers on an online website comprised of those who love amateur chemistry and science suggest, you are in an extremely un-enviable position, which may not be due to your own fault, and I would highly recommend examining how to improve your situation.

I do mean this as an attempt to be helpful and not just make light of your situation.

[Edited on 15-5-2015 by Chemosynthesis]

blogfast25 - 15-5-2015 at 10:22

Ionophores? Blimey, occasionally one does learn something at SM!

DistractionGrating - 15-5-2015 at 10:56

Checking solubility tables, I see that Calcium molybdate is only sparingly soluble (0.004099 g / 100 mL), while Magnesium molybdate is much more soluble (13.7 g / 100 mL). Now, I know nothing about working with molybdates, but would it be possible to add magnesium molybdate in order to precipitate out calcium molybdate? If so, then of course, it would be helpful to know how much calcium is present in order to know how much magnesium molybdate to add.

[Edited on 15-5-2015 by DistractionGrating]

blogfast25 - 15-5-2015 at 11:14

Quote: Originally posted by DistractionGrating

Checking solubility tables, I see that Calcium molybdate is only sparingly soluble (0.004099 g / 100 mL), while Magnesium molybdate is much more soluble (13.7 g / 100 mL). Now, I know nothing about working with molybdates, but would it be possible to add magnesium molybdate in order to precipitate out calcium molybdate? If so, then of course, it would be helpful to know how much calcium is present in order to know how much magnesium molybdate to add.

[Edited on 15-5-2015 by DistractionGrating]

The cost of the molybdate would not justify trying to 'save' that MgCl2. I'd bet good money on that being the case.

MrHomeScientist - 15-5-2015 at 12:27

It always amuses me when people specify they want the "best and easiest" way to do something. Really? I never would have guessed.

Unfortunately I only know the worst and hardest ways.

DistractionGrating - 15-5-2015 at 12:41

Quote: Originally posted by blogfast25

The cost of the molybdate would not justify trying to 'save' that MgCl2. I'd bet good money on that being the case.

Fair enough, but does my idea have merit? Just trying to learn.

blogfast25 - 15-5-2015 at 14:25

Quote: Originally posted by DistractionGrating

Fair enough, but does my idea have merit? Just trying to learn.

Assuming your solubility data are correct, then yes.

Eli25 - 15-5-2015 at 20:40

Hiiiiii.... It was so nice seeing your comments...Actually I'm here to learn more and as we say in our country " every mind has a new thought", I'm looking forward to gain more different ideas and suggestions....I'm not here to find the exact answer for those who said it's not our problem.....

Believe me your comments help me consider all different ways to solve my problem and I'm learning alot. I do appreciate your help.

Eli25 - 15-5-2015 at 20:46

by the way where can I find the rightest values of solubility and Ksp?

I want to know if CaSo4 is more insoluble or CaCo3? to know weather adding MgCo3 works better or adding MgSo4 to my solution to gain purer MgCl2!!!!

DistractionGrating - 15-5-2015 at 20:59

In pure water, absent any other ions, CaSO4 is more soluble than CaCO3, but solubility is affected by the presence of other ions. For example, CaCO3 is more soluble in seawater than it is in freshwater. Then, there is the "common ion effect" that can make some substances much less soluble when a common cation or anion is present. I'm still trying to figure this all out myself.

Chemosynthesis - 15-5-2015 at 21:43

Quote: Originally posted by Eli25

by the way where can I find the rightest values of solubility and Ksp?

I want to know if CaSo4 is more insoluble or CaCo3? to know weather adding MgCo3 works better or adding MgSo4 to my solution to gain purer MgCl2!!!!

If handbooks like CRC, Lange's and Merck don't suit you, consider checking what are known as critical tables.

blogfast25 - 16-5-2015 at 05:43

Quote: Originally posted by Eli25

Solubility products:

http://bilbo.chm.uri.edu/CHM112/tables/KspTable.htm

Quote: Originally posted by DistractionGrating

Then, there is the "common ion effect" that can make some substances much less soluble when a common cation or anion is present. I'm still trying to figure this all out myself.

It is, simply put, the consequence of solubility products:

AB(s) < === > A+(aq) + B-(aq)

Ks = [A+] x [B-]

If the solution now also contains A+(aq) OR B-(aq) coming from another soluble compound, then it will push the equilibrium further to the left, rendering the compound AB even more insoluble. Example: CaCO3 + K2CO3. Discuss.

[Edited on 16-5-2015 by blogfast25]

DistractionGrating - 16-5-2015 at 09:04

I have a basic grasp of the common ion effect. The part I'm wrestling with is how, for example, the presence of Mg and/or SO4 can make CaCO3 *more* soluble, as is the case in seawater.

blogfast25 - 16-5-2015 at 09:42

Quote: Originally posted by DistractionGrating

I have a basic grasp of the common ion effect. The part I'm wrestling with is how, for example, the presence of Mg and/or SO4 can make CaCO3 *more* soluble, as is the case in seawater.

One way the presence of 'spectator ions' can affect solubility is that they influence the so-called ionic strength of a solution. This has in turn an effect on the activity coefficients of the solvated species.

CaCO3's solubility is also, for obvious reasons, quite pH dependent. Dissolved CO2 also has an effect (formation of bicarbonates).

See Debye - Huckel theory for a better understanding. Oh, and good luck: it's complicated theory and that's one of the reasons we mostly assume the activity coefficients are one (which is only strictly true in infinite dilutions)

Edit: an interesting study that takes activity coefficients into account in the determination of the solubility product of Malachite, here:

http://www.nrcresearchpress.com/doi/pdf/10.1139/v57-177

[Edited on 16-5-2015 by blogfast25]